2 1.1 Chemistry—The Science of Everyday Experience Chemistry is the study of matter—its composition,properties, and transformations.Matter is anything that has mass and takes up volume.Matter can be:Naturally occurring:cottonsanddigoxin, a cardiac drugSynthetic (human-made):nylonStyrofoamibuprofen
3 1.2 States of Matter The Solid State: A solid has a definite volume. It maintains its shape regardless of its container.Solid particles lie close together in a regular pattern.
4 1.2 States of Matter The Liquid State: A liquid has a definite volume. It takes the shape of its container.Liquid particles are close together but can move past one another.
5 1.2 States of Matter The Gas State: A gas has no definite shape; it assumes the shape of its container.It has no definite volume; it assumes the volume of its container.Gas particles are very far apart and move around randomly.
6 1.2 States of Matter Physical properties can be observed or measured without changing the composition of the material.boiling pointmelting pointsolubilitycolorodorstate of matter
7 1.2 States of MatterA physical change alters the material without changingits composition (changes in state).
8 1.2 States of Matter Chemical properties determine how a substance can be converted into another substance.Chemical change is the chemical reaction thatconverts one substance into another (Chapters 5 and 6).
9 1.3 Classification of Matter All matter can be classified as either a pure substance or a mixture.I. Pure SubstancesA pure substance is composed of only a single component (atom or molecule).It has a constant composition, regardless of sample size or origin of sample.It cannot be broken down to other pure substances by a physical change.
10 1.3 Classification of Matter All matter can be classified as either a pure substance or a mixture.I. Pure SubstancesTable sugar (C12H22O11) and water (H2O) are bothpure substances:
11 1.3 Classification of Matter All matter can be classified as either a pure substance or a mixture.II. MixturesMixtures are composed of more than one component.They can have varying composition (any combination of solid, liquid, and gas).Mixtures can be separated into their components by a physical process.
12 1.3 Classification of Matter All matter can be classified as either a pure substance or a mixture.II. MixturesSugar dissolved in water is a mixture.
13 1.3 Classification of Matter A pure substance is classified as an element or a compound.I. An element is a pure substance that cannot be broken down by a chemical change.aluminum metal (Al)
14 1.3 Classification of Matter A pure substance is classified as an element or a compound.II. A compound is a pure substance formed by chemically joining two or more elements.table salt (NaCl)
16 1.4 Measurement Every measurement is composed of a number and a unit. The number is meaningless without the unit.Examples:proper aspirin dosage = 325 (milligrams or pounds?)a fast time for the 100-meter dash = (seconds or days?)
17 1.4 Measurement A. The Metric System Each type of measurement has a base unit in the metric system..
18 1.4 Measurement A. The Metric System Other units are related to the base unit by a power of 10.The prefix of the unit name indicates if the unit is larger or smaller than the base unit.
19 1.4 Measurement B. Measuring Length The base unit of length is the meter (m).1 kilometer (km) = 1,000 meters (m)1 km = 1,000 m1 millimeter (mm) = meters (m)1 mm = m1 centimeter (cm) = 0.01 meters (m)1 cm = 0.01 m
20 1.4 Measurement C. Measuring Mass Mass is a measure of the amount of matter in an object.Weight is the force that matter feels due to gravity.The base unit of mass is the gram (g).1 kilogram (kg) = 1,000 grams (g)1 kg = 1,000 g1 milligram (mg) = grams (g)1 mg = g
21 1.4 Measurement D. Measuring Volume The base unit of volume is the liter (L).1 kiloliter (kL) = 1,000 liters (L)1 kL = 1,000 L1 milliliter (mL) = liters (L)1 mL = LVolume = Length x Width x Height= cm x cm x cm= cm31 mL = 1 cm3 = 1 cc
23 1.5 Significant FiguresAn exact number results from counting objects or ispart of a definition.10 fingers10 toes1 meter = 100 centimetersAn inexact number results from a measurement orobservation and contains some uncertainty.15.3 cmgmL
24 1.5 Significant Figures A. Determining Significant Figures Significant figures are all the digits in a measurednumber including one estimated digit.All nonzero digits are always significant.65.2 g65.2 ggg3 sig. figures6 sig. figures
25 1.5 Significant Figures A. Determining Significant Figures Rules for Zero:Rule 1: A zero counts as a significant figure whenit occurs:between two nonzero digits29.05 g29.05 gmLmL4 sig. figures5 sig. figuresat the end of a number with a decimal placecmcm620. lb620. lb5 sig. figures3 sig. figures
26 1.5 Significant Figures A. Determining Significant Figures Rules for Zero:Rule 2: A zero does not count as a significant figurewhen it occurs:at the beginning of a numbermgmg0.008 mL0.008 mL3 sig. figures1 sig. figureat the end of a number that does not have a decimal2570 m2570 mmm3 sig. figures5 sig. figures
27 1.5 Significant Figures B. Rules for Multiplication and Division The answer has the same number of significant figuresas the original number with the fewest significant figures.4 sig. figures351.2 miles351.2 milesmiles=hour5.5 hour5.5 hour2 sig. figuresAnswer must have2 sig. figures.
28 1.5 Significant Figures B. Rules for Multiplication and Division to be retainedto be droppedmiles=64 mileshourhourfirst digit to be dropped2 sig. figuresAnswerIf the first digitto be dropped is:Then:between 0 and 4drop it and all remaining digitsbetween 5 and 9round up the last digitto be retained by adding 1
29 1.5 Significant Figures B. Rules for Multiplication and Division
30 1.5 Significant Figures C. Rules for Addition and Subtraction The answer has the same number of decimal placesas the original number with the fewest decimal places.10.11 kg10.11 kg2 decimal places3.6 kg3.6 kg1 decimal place6.51 kganswer must have1 decimal placefinal answer1 decimal place=6.5 kg
31 In scientific notation, a number is written as: y x 10xy x 10xExponent:Any positive or negativewhole number.Coefficient:A number between1 and 10.
32 1.6 Scientific Notation 2,500 0.036 2500 0.036 2.5 x 103 3.6 x 10−2 HOW TO Convert a Standard Number to Scientific NotationExampleConvert these numbers to scientific notation.2,5000.036Move the decimal point to give a numberbetween 1 and 10.Step 25000.036Multiply the result by 10x, wherex = number of places the decimal was moved.Step move decimal left,x is positivemove decimal right,x is negative2.5 x 1033.6 x 10−2
33 Converting a Number in Scientific Notation to a Standard NumberWhen the exponent x is positive, move the decimal point x places to the right.2.800 x 102 =280.0When the exponent x is negative, move the decimal point x places to the left.2.80 x 10–2 =0.0280
34 1.7 Using the Factor-Label Method A. Conversion Factors Conversion factor: A term that converts a quantity inone unit to a quantity in another unit.originalquantityconversion factordesiredx=Conversion factors are usually written as equalities.2.21 lb = 1 kgTo use them, they must be written as fractions.2.21 lb1 kgor
35 units are treated like numbers make sure all unwanted units cancel 1.7 Using the Factor-Label Method B. Solving a Problem Using One Conversion FactorFactor-label method: Using conversion factors to convert a quantity in one unit to a quantity inanother unit.units are treated like numbersmake sure all unwanted units cancelTo convert 130 lb into kilograms:130 lbxconversion factor? kg=original quantitydesired quantity
36 1. 7 Using the Factor-Label Method B 1.7 Using the Factor-Label Method B. Solving a Problem Using One Conversion Factor2.21 lb1 kgAnswer2 sig. figures130 lbxor1 kg2.21 lb=59 kgThe bottom conversion factor hasthe original unit in the denominator.The unwanted unit lb cancels.The desired unit kg does not cancel.
37 1.7 Using the Factor-Label Method HOW TO Solve a Problem Using Conversion FactorsHow many grams of aspirin are in a 325-mgtablet?ExampleIdentify the original quantity and the desiredquantity, including units.Step original quantitydesired quantity325 mg? g
38 1.7 Using the Factor-Label Method HOW TO Solve a Problem Using Conversion FactorsWrite out the conversion factor(s) neededto solve the problem.Step 1 g = 1000 mgThis can be written as two possible fractions:1000 mg1g1 g1000 mgorChoose this factor tocancel the unwantedunit, mg.
39 1.7 Using the Factor-Label Method HOW TO Solve a Problem Using Conversion FactorsStep Set up and solve the problem.1 g1000 mg0.325 g0.325 g325 mg325 mgx=3 sig. figures3 sig. figuresUnwanted unitcancelsWrite the answer with the correct numberof significant figures.Step 
40 1. 7 Using the Factor-Label Method C 1.7 Using the Factor-Label Method C. Solving a Problem Using Two or More Conversion FactorsAlways arrange the factors so that the denominator inone term cancels the numerator in the preceding term.How many liters is in 1.0 pint?1.0 pintoriginal quantity? Ldesired quantityTwo conversion factors are needed:2 pints = 1 quart1.06 quarts = 1 liter2 pt1 qt1 qt2 pt1.06 qt1 L1 L1.06 qtororFirst, cancel pt.Then, cancel qt.
41 Set up the problem and solve: 1.7 Using the Factor-Label Method C. Solving a Problem Using Two or More Conversion FactorsSet up the problem and solve:1 qt2 pt1 L1.06 qt1.0 pt1.0 ptxx=0.47 LL2 sig. figures2 sig. figuresWrite the answer with the correct number ofsignificant figures.
42 To convert from oC to oF: To convert from oF to oC: 1.9 TemperatureTemperature is a measure of how hot or cold an object is.Three temperature scales are used:Degrees Fahrenheit (oF)Degrees Celsius (oC)Kelvin (K)To convert from oC to oF:To convert from oF to oC:oC = oF − 321.8oF = 1.8(oC)To convert from oC to K:To convert from K to oC:K = oCoC = K − 273
43 1.9 Temperature Comparing the Three Temperature Scales
44 1.10 Density and Specific Gravity A. Density Density: A physical property that relates the mass ofa substance to its volume.mass (g)density=volume (mL or cc)To convert volume (mL) to mass (g):To convert mass (g) to volume (mL):gmLmLx=ggx=mLmLginverse of densitydensity
45 1.10 Density and Specific Gravity A. Density Example:If the density of acetic acid is 1.05 g/mL, what isthe volume of 5.0 grams of acetic acid?5.0 g? mLoriginal quantitydesired quantityDensity is the conversion factor, and can bewritten two ways:1.05 g1 mL1 mL1.05 gChoose the inverse densityto cancel the unwanted unit, g.
46 1.10 Density and Specific Gravity A. Density Set up and solve the problem:1 mL1.05 g5.0 g5.0 gx=mL4.8 mL2 sig. figures2 sig. figuresUnwanted unitcancelsWrite the final answer with the correct numberof significant figures.
47 1.10 Density and Specific Gravity B. Specific Gravity Specific gravity: A quantity that compares the densityof a substance with the density of water at thesame temperature.density of a substance (g/mL)density of water (g/mL)specific gravity=The units of the numerator (g/mL) cancel theunits of the denominator (g/mL).The specific gravity of a substance is equal to itsdensity, but contains no units.